Molten carbonate fuel cell research

Electrochimica Acta 45 (1999) 1025 – 1037 www.elsevier.nl/locate/electacta

Molten carbonate fuel cell research Part I. Comparing cathodic oxygen reduction in lithium/potassium and lithium/sodium carbonate melts Kosmas Janowitz, Michael Kah, Hartmut Wendt * Institute for Chemical Technology, TU-Darmstadt, Petersenstraße 20, D-64287 Darmstadt, Germany Received 3 February 1999; received in revised form 16 July 1999

Abstract This paper reports on microkinetic measurements of cathodic oxygen reduction at immersed flag electrodes in Li/K carbonate melts and macrokinetic measurements of the same reaction in laboratory fuel cells with eutectic Li/K and Li/Na carbonate electrolyte. The microkinetic measurements confirm that at 650°C oxygen reduction is a relatively fast reaction with exchange current densities of the order of 10 − 2 A cm − 2 at gold, lithiated nickel oxide and lithium cobaltite electrodes, so that charge transfer hindrance might be of little relevance for the macrokinetics of molten carbonate fuel cell (MCFC) cathodes. O2 reduction seems to be faster on Lithium cobaltate than on lithiated nickel oxide. The results on gold electrodes indicate that in Li/K carbonate melts the species which transports oxygen may be peroxocarbonate. Macrokinetic measurements show that mass transfer of oxygen determines current – voltage curves, in particular at low oxygen partial pressures but also at low carbon dioxide partial pressures of the order of 10 − 1 bar and less, which are relevant to the technical operating conditions of MCFCs. CO2 transport bears sizeably on the cathodic current–voltage correlation at low partial pressures. Since the oxygen solubility is lower in Li/Na than in Li/K carbonate melts, oxygen overpotentials become much higher with lean gas mixtures at higher current densities with Li/Na carbonate melts. Therefore in the practical sense the Li/K carbonate melt seems to be overwhelmingly superior to the Li/Na carbonate eutectic in MCFC technology. If the morphology of MCFC cathodes can be significantly improved to remove oxygen and carbon dioxide mass transfer hindrance, or if pressurised cells are used in which O2 and CO2 partial pressures exceed 0.25 bar even for lean gases, the less volatile and more conductive the Li/Na carbonate melt will be a superior alternative to the Li/K carbonate eutectic as an MCFC electrolyte. © 1999 Elsevier Science Ltd. All rights reserved. Keywords: Fuel cells; MCFC; Alkali carbonate melts; Cathodic oxygen reduction

1. Introduction Since the beginning of molten carbonate fuel cell (MCFC) development the most effective electrolyte for * Corresponding author. Tel.: +49-6151-162-265; fax: + 49-6151-164-788. E-mail address: [email protected] (H. Wendt)

this energy-converter had been the subject of debate. In early works the ternary Li/Na/K carbonate eutectic was used [1], but for more than 20 years the Li/K carbonate eutectic with 68/32 mol/mol lithium carbonate/potassium carbonate has been the electrolyte of choice for MCFCs operating at 650°C. Recently the use of this electrolyte was discussed by one of the present authors (H.W.), as its potassium volatility (in the form of KOH) seems to be too high for safe long-term MCFC

0013-4686/99/$ - see front matter © 1999 Elsevier Science Ltd. All rights reserved. PII: S 0 0 1 3 - 4 6 8 6 ( 9 9 ) 0 0 3 0 5 - 9

1026

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

operation [2]. The electrolytic conductivity of the Li/Na carbonate is better by a factor of roughly two. Since the volatility of sodium in Li/Na carbonate eutectic is lower by a factor of 10 than that of potassium above Li/K eutectic, this electrolyte was considered to be a better choice by several other authors [3–5]. Polarisation measurements with so-called ‘fat’ cathode gas, which contains 33 vol. % O2 and 66 vol. % CO2 obviously indicated that the Li/Na carbonate electrolyte was kinetically adequate, although its oxygen solubility is known to be by a factor of two lower than that in Li/K carbonate eutectic [6]. However, fuel cell developers reported unforeseen difficulties with this electrolyte, since they obtained much lower power densities than predicted using lean cathode gas compositions, which today are the state of the art for MCFC operation. This was the reason for the present authors to reinvestigate the electrode kinetics of cathodic oxygen reduction in both electrolytes. The kinetic investigation was performed on two different levels: 1. A microkinetic investigation confined to non-steady and steady-state measurements at smooth immersed flag electrodes in lithium/potassium carbonate eutectic to determine exchange current densities at three different electrode materials (Au, NiO and LiCoO2) and to obtain additional information on the nature of transported oxygen species. 2. A macrokinetic measurement to determine current– voltage correlations of real fuel cells with oxygen cathodes whose nickel oxide electrodes possess the same porosity, morphology and composition as industrial MCFC cathodes, since they were produced from the same precursors and according to the same procedures as truly-technical MCFC cathodes. The aim was to determine formal reaction orders at these electrodes from the charge transfer resistances for fat and for lean gas compositions in operating fuel cells to obtain a reliable judgement as to whether Li/Na carbonate eutectic may or may not be used in MCFCs. The purposes were to obtain electrode kinetic data with respect to exchange current densities and reaction orders for carbon dioxide and oxygen which can be compared with the ample literature on the subject and to correlate these data with macrokinetic information from current density–overpotential curves of fuel cell cathodes. It was expected from literature data that the exchange current densities at oxygen and carbon dioxide pressures of between 0.1 and 0.5 bar were of the order of magnitude of 10 − 2 A cm − 2 which means standard rate constants are of the order of 10 − 1 cm s − 1 (assuming reactant concentrations in the melt of 10 − 7 to 10 − 6 mol cm − 3). Electrode kinetic steadystate measurements should therefore be performed using methods which allow for mass transfer coefficients

of higher than 10 − 1 cm s − 1, which simply do not exist. Therefore very fast pulse methods had been applied which gave access to measuring times of B1 ms. Such measurements must, however, cope with the difficulty of separating double layer charging currents from Faradaic currents. Electrode capacitancies at the metal or metal oxide/molten salt electrolyte interface are relatively high, measuring ca. 100 mF cm − 2. In contrast electrode/aqueous electrolyte interfaces have capacitances of a few tens of mF cm − 2. Therefore, the rise time of the pulse method had to be between 10 and 20 ms. Three different fast measurements were compared for determining exchange current densities: impedance spectroscopy (IS) coulostatic pulsing (CP) and ‘potentiostatic pulsing’ (PP). Impedance spectroscopy, in which Nyquist plots are evaluated by determining slopes of the plots at the highest frequencies were shown to be insufficiently accurate. For instance, the i0 values obtained at gold electrodes at 650°C in O2/ CO2 = 9:1 atmosphere varied between 20 and greater than 250 mA cm − 2, and were thus unreliable. In coulostatic pulsing the electrode potential is raised by charging the electrode in microseconds by a capacitor (for instance of several nF) charged to several volts (1 – 10 V depending on the intended change of the electrode potential). This method which was described by Delahay [7] and later by van Leuwen [8] and was used by Uchida and his group [9 – 11]. It turned out to have a poorer time resolution in practice than ‘potentiostatic pulsing’. The latter method does not really allow the application of an infinitely short voltage jump to the electrode. Fig. 1(a) gives in the left insert the schematic of RC combinations embodied in the cell by the working electrode with charge transfer resistor and double layer capacitance and the electrolyte resistance between the reference and the working electrode. The points A, B and C mean counter, reference and working electrodes. Double layer charging typically needs a rise-time which is determined by the electrolyte resistance (perhaps 0.1 – 0.5 V cm2) and by the double layer capacitance(perhaps 100 mF cm − 2) of 5× 10 − 5 to several 10 − 4 s. The electrode potential is adjusted to the equilibrium potential at the beginning of the experiment. In Fig. 1(a) the very fast linear voltage sweep (instead of a true step) is shown. This is completed after 20 ms and the electrode potential changes after the short period of the voltage sweep according to the exponential function Eq. (1) h(t)= h (1− exp{− t/t})

(1)

with t=

1 RelCD

(2)

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

1027

Fig. 1. Schematic presentation of (a) increase of control potential with time and corresponding change of overpotential at the working electrode, plotted vs. square root of time. Inserts: left; schematic of equivalent for electrolyte resistance, charge transfer resistance and double layer capacitance, and right: the rapid increase in total current in response to the change of control potential, with current peak at the end of the control potential step, followed by exponential decay of the charging current and an exponential approach of overpotential towards the externally applied overpotential. The Faraday or charge transfer current increases slowly to a flat maximum and then begins to decay according to t − 0.5. (b) Enlarged fraction of the insert plot in (a) (right hand side) showing the increase of charge transfer current density with a peak at t 0.5 equal to approx. 0.2 ms0.5, and linear decay of charge transfer current with the square root of time. The current densities at zero time is obtained by extrapolating the charge transfer current to t= t/2. This figure also shows the decay of the charging currents, ic, with time.

The potential and current density are plotted versus the square root of the time after triggering the voltage step. In the insert of Fig. 1(a,b) the different partial currents which are flowing through the cell in response to these potential changes are recorded. After triggering the commanding potential sweep the current rises steeply due to double layer charging, and peaks at the moment the fast linear potential sweep is halted at the

predetermined value DE . At this moment double layer charging is far from completion, and the electrode overpotential is only 10 – 20% of the electrode overpotential eventually established in the long term. Charging the double layer occurs from there onwards with exponentially decaying charging current according to Eq. (3), while the overpotential increases according to Eq. (1)

1028

 !

ic = imax exp −

t −t Rel·Cd

"

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

(3)

If during double layer charging a steadily increasing Faradaic current flows, then the overpotential changes due to charging according to Eq. (3) are given by Eqs. (4a) and (4b) DE−IcRel

h=



1+i0ne fRel 1−

2H

p

t

n

(4a)

with H=

i0exp{aFh/RT} neFCOx DOX

(4b)

The Faradaic or charge transfer current if or ict, respectively, is almost negligible compared to the charging ic current at the beginning of double layer charging. It then increases with growing overpotential, and later begins to decrease due to mass transfer hindrance and growing depletion of depolariser ‘Ox’ at the cathode surface. Correspondingly, the Faradaic current decreases according to the t − 0.5 law. In addition to pulse methods, steady-state measurements of cathodic oxygen reduction at immersed flag electrodes under conditions of well defined free thermal convection have been used to obtain information on the transport parameter cox. The same purpose served cyclic voltammetry at sufficiently fast sweep rates to avoid interference of free convection. These measurements yield information on coxD 0.5, i.e. also on cox. Both measurements were used to obtain formal O2 and CO2 reaction orders for the formation of the species formed by dissolution of oxygen in the melt.

2. Experimental For electrode kinetic investigations with immersed flag electrodes, almost identical equipment to that of Nicholson and Appleby in their well-known early kinetic experiments [12–14] was used. Three porous alumina plates, of 5 cm thickness each, above the cell served as heat insulators against the lid of the closed alumina tube, which contained the alumina crucible with the melt and electrodes. The 1 ×1 cm2 gold flag electrode connected to the 0.15 mm inner diameter supporting alumina tube by a gold plug was fully immersed in the electrolyte to avoid difficulties with the meniscus [15,16], i.e. the meniscus developed on the supporting alumina tube and not on the electrode surface. Since it is impossible to attach a Luggin capillary to the oxygen reference electrode, the non-negligible IR drops must be accounted for by determining them either by current interruption and/or by analysing impedance spectra. Kinetic measurements were performed on gold, nickel oxide and lithium cobaltite

(LiCoO2) electrodes. The latter two oxide ceramic electrodes were prepared on both sides of the gold flag electrode on which different amounts of nickel and cobalt had been previously deposited from aqueous galvanic baths. The nickel bath contained 240 g/l NiSO4, 45 g/l NiCl2 and 40 g/l H3BO4 and the deposition current density was 6 mA cm − 2 at 50°C. The cobalt bath contained 450 g/l CoSO4, 20 g/l CoCl2, 40 g/l H3BO4, 20 g/l KCl [17], and the deposition was performed at 4 mA cm − 2 at 50°C. Both metal deposits were even and smooth. The respective oxides were prepared in-situ by chemical oxidation with dissolved oxygen in the gas-stirred electrolyte, with monitoring of the electrode potential using a method previously applied by Tomczyk and colleagues [18,19]. Macrokinetic measurements were performed in a typical apparatus with horizontal cells of 3 cm2 geometric surface area, squeezed between two vertical alumina tubes of 2 cm inner diameter and half a meter length, which were described in a review by Selman and Marianowski [20]. The cells were constructed from foil precursors as described elsewhere [21] by one of the authors and his colleagues, and were formed in-situ by careful heating with introduction of the melt either by inclusion of an electrolyte foil between matrix and cathode foil, or by supplying it to the circular rim of the cell. The electrolyte loading amounted to 0.3 g cm − 2. This loading is the approximate optimal load for maximal power density. The reference electrode was placed outside the cell at an extension of the wet seal, which sat on a tubular alumina collar. It consisted of a 3 mm ¥ spiralled 0.2 mm diameter gold wire pressed against the electrolyte-soaked matrix material. This reference was supplied with a gas mixture of 33/67 vol./ vol. oxygen/carbon dioxide via an alumina tube of 4 mm inner diameter. The flat cell was placed in the middle of a 40 cm long upright tubular furnace. This was thermostatted by regulated heating to a precision of 9 2°C. The anode gas always contained 70 vol.% hydrogen and 18 vol.% carbon dioxide, together with 12 vol.% water vapour. The anode and cathode gases were never utilised to more than 10 %. Standard operating temperature was 650°C, but calibrated cell temperatures as low as 550° and as high as 750°C could be used to measure activation energies. In the experiments, the reference, counter and working electrodes of the cell were connected to an Amel Instruments 2049 potentiostat, and the electrode potential was adjusted to the equilibrium potential at the beginning of the kinetic experiment. For PP experiments a function generator of our own design, with a step function rise time of 20 ms was coupled to the input of the potentiostat and the cell current measured by a probe resistor. The cell was monitored with a Vukodigitalscope 22-16 which had a sampling frequency of 2 MHZ. Data storage and processing were performed

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

1029

Fig. 2. Charge transfer current density vs. t 0.5 obtained from total currents by subtracting the charging currents. The charging currents for three different overpotentials ( − 30, − 50 and −90 mV) are scanned by three independent measurements.

using a 286 IBM PC unit. For linear sweep voltametry a Philips function generator was used. In-cell experiments were performed by increasing the cathodic currents in 20 mA cm − 2 steps. A time period of 15 min. was always allowed to elapse before reading the cathode potential.

3. Results

3.1. Microkinetic measurements 3.1.1. Fast methods for determining exchange current densities 3.1.1.1. The gold electrode. Fig. 2 shows three typical recordings of the Faradaic current transients in PP measurements obtained on gold flag electrodes immersed in a stationary Li2CO3/K2CO3 eutectic melt (62/38 mol/mol) saturated at 650°C with O2/CO2 at 0.9 bar/0.1 bar for three different potential steps (−30, − 50 and −90 mV). The measurements were repeated three times each and the points depicted were obtained by subtracting the capacitive charging currents from the total current transient. These currents are also plotted in Fig. 2. The figure allows an estimate of reproducibility and scatter. The scatter is greatest in the ascending part during double layer charging. The requirement is to obtain the true current density at ‘virtual zero’

time, which is not identical with the beginning of the pulse. In effect, there is no time ‘zero’ in such PP experiments, since the Faradaic current begins to flow and the oxidised species begins depleting as soon as the elctrode potential begins to deviate from the equilibrium potential. From Fig. 1(a,b) it is clear that from t − 0.5 = 0.2 ms − 0.5 (t = 40 ms) to 0.45 ms0.5 (t = 200 ms) the Faradaic current increases up to the peak and then begins to decrease according to t − 0.5. It would be a practical compromise to choose as the ‘virtual zero’ time from the t 0.5 scale, when the concentration of the Ox species near the electrode begins to deplete the arithmetic mean value between the onset of the Faradaic current, and that for the t 0.5 value at which the Faradaic current peaks. This value is somewhat different for the three different curves but is approximately 0.32 90.03 ms0.5. The three straight lines projecting towards the current density axis allow extrapolation of current densities at virtually zero time equal to 25, 16 and 9 mA cm − 2 for overpotentials of −90, −50 and −30 mV. From an Allen – Hickling plot, ln{it = 0/[exp F2h/RT −1]} versus h, one obtains the exchange current density i0 = 13.690.2 mA cm − 2

(5)

[650°, Au, O2/CO2 = 0.9 bar/0.1 bar] The double layer capacitance and electrolyte resistance obtained from the time transient of the charging current are 135 mF cm − 2 and 0.41 V cm − 2 and the cathodic charge transfer coefficient is ac = 0.62.

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

1030

2 C LiCoO : 300920 mF cm − 2 D

Fig. 3. Extrapolation of exchange current densities for oxygen reduction at lithiated NiO and LiCoO2 cathodes vs. oxide layer thickness. At oxide layer thickness equal to 0.25 mm, the exchange current density of planar oxide electrodes is assumed to be approached. The extrapolation for LiCoO2 shows that the support material, Au, does not ‘shine through’ the porous oxide surface.

3.1.1.2. Nickel oxide and lithium cobaltite electrodes. An intrinsic problem in investigating electrode kinetics at oxide ceramic electrodes which usually are porous is determining the true surface area for calculation of true current densities. For porous nickel or lithium cobaltite electrodes, i0 values for oxide coatings of variable thickness were determined. Extrapolation of the measured effective current densities down to the thickness which approaches the mean crystallite or particle diameter of the oxide ceramic material would give a fair approximation to the current density of a smooth oxide electrode, provided the porosity of the oxide coating is 50 % or less. The porosity of these electrodes obtained by in-situ oxidation is approximately 0.3, and the mean particle diameter approaches 0.25–0.3 mm. Fig. 3 shows for the current densities of nickel oxide coated gold electrodes with coatings of 0.25, 0.5 and 0.75 mm thickness. One obtains by extrapolation to 0.25 mm thickness i NiO :i(0.25 mm)= 6.390.7 mA cm − 2 0

(6a)

with an electrode capacitance of −2 C NiO D :200 915 mF cm

i

:i(0.25 mm)=139 2 mA cm

with the specific capacitance

with the transfer coefficient a =0.579 0.03. It should be mentioned that the coulostatic pulse method yields current densities at Au, NiO and LiCoO2 which are approximately 15 – 20 % higher. Table 1 compares i0, CD and a values for oxygen reduction with O2/CO2, 0.9 bar/0.1 bar atmosphere in Li/K carbonate eutectic for the three different electrode materials. The results will be compared with literature data in Section 4. However, the linear dependencies of current density on oxide layer thickness itself indicates that the material is porous. That the i0 for NiO is lower than for Au is also an indicator that the gold support does not ‘shine through’ in the kinetic measurements with a porous nickel oxide layer. From the Arrhenius plot log i0 versus 1/T, an activation energy of E*= 5192 kJ mol − 1 is obtained for the gold electrode. At the gold electrode reaction orders were determined for oxygen as: nO2 = 0.3390.1

(8a)

and for carbon dioxide as: nCO2 = 0.1390.1

(8b)

3.1.2. Slow methods 3.1.2.1. Steady-state current –6oltage cur6es. Natural convection is often undesirable at electrodes immersed in molten salts, but is usually significant since the so-called Rayleigh number Ra is high. Ra is the characteristic dimensionless quantity for free thermal convection in fluids contained in vessels of dimensions in which small spatial temperature differences (DT) prevail and induce natural convection which even might become turbulent. The Rayleigh number is given by: Ra= GrPr=(r 4gbDT/ln 2)(n/a)

(9)

where r= radius of the vessel, h= height of the vessel, g= gravitational constant, DT= temperature difference from top to bottom of the vessel, b =thermal expansion, n= kinematic viscosiy and a=temperature conductivity of the melt. Ra is of the order of 103 for our alumina crucible. Table 1 Electrode kinetic data of cathodic oxygen reduction at three different electrode materials in Li2CO3/K2CO3 (62/38 mol/ mol) eutectic at 650°C saturated with O2/CO2 =0.9 bar/0.1 bar

(6b)

a was 0.590.005. Fig. 3 explicates the same procedure also for lithium cobaltite with the result LiCoO2 0

(7b)

−2

(7a)

Electrode material

i0 (mA cm−2)

CD (mF cm−2)

a

Au NiO LiCoO2

13.6 90.2 6.3 9 0.7 13 9 2

135 910 200 915 300 920

0.62 0.5 0.57

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

1031

Fig. 4. Current – voltage curves for cathodic oxygen reduction at real MCFC cathodes at constant partial pressures of CO2 of 0.13 bar and variable oxygen partial pressure. The dotted lines were measured with lithium/sodium carbonate melt instead of Li/K carbonate melts.

Lu and Selman [22] obtained typical steady-state current–voltage curves at the gold flag electrode in the Li2CO3/K2CO3 eutectic melt due to thermal convection. These steady-state investigations of oxygen reduction were performed under non-pressurised atmospheres at a constant 0.3 bar CO2, but at five different oxygen partial pressures (0.02, 0.04, 0.1, 0.2 and 0.4 bar, balance nitrogen). All current–voltage curves showed more or less well defined two different cathodic waves and mass-transfer-limited current densities at overpotentials from −100 to −160 mV, as described by Appleby and Nicholson. The meaning or explanation of the first wave, which is well developed at the highest oxygen pressure at 0 to −70 mV overpotential at oxygen partial pressures as low as 0.2 bar was not investigated in detail. Appleby and Nicholson tried to attribute the two different cathodic processes to the reduction of peroxide and superoxide anions. Another explanation may be the reduction of peroxide anions with and without participation of carbon dioxide. We concentrate here on the interpretation of the mass transfer limited current densities at potentials more negative than −120 to 160 mV, i.e. for the sum of both waves. Plotting the logarithm of the sum of the two waves (i.e. mass transfer limited current densities) versus the logarithm of the oxygen pressure gives a reaction order for oxygen species transported in the melt. The reaction order is 0.5. The mass transfer reaction order of carbon dioxide is low, and cannot easily be exactly determined. In the carbon dioxide partial pressure range from 0.03 to 0.1 bar, it is −0.279 0.15 and between 0.1 and 0.7 bar it is slightly positive (+0.12 9

0.08). Hence, nCO2 is close to zero in the CO2 partial pressure range from 0.03 to 0.7 bar, and the mass transfer limited current density of oxygen is scarcely influenced by the carbon dioxide partial pressure.

3.1.2.2. Linear scan 6oltammetry. With cyclic voltammetric measurements peak current densities are evaluated which are determined by the transport quantity cOXD 1/2 OX. The method was applied by Appleby and Nicholson for the first time, and was repeated by Uchida and colleagues [12 – 14], Selman [22], and more recently by Appleby and co-workers [23 – 25]. We therefore dispense with a detailed description of our results, and report only the formal reaction orders for O2 and CO2. In good agreement with data from mass transfer limited current densities, we obtained from linear sweep voltammetry nO2(LSV) = 0.479 0.04

(10a)

and nCO2(LSV) = 0.059 0.05

(10b)

3.1.3. Macrokinetic measurements from current –6oltage plots in real fuel cells 3.1.3.1. O2 reduction kinetics in Li2CO2 /K2CO3 melts. Fig. 4 shows IR-corrected cathodic overpotential current density correlations measured in cells with Li2CO3/ K2CO3 eutectic electrolyte at a constant carbon dioxide partial pressure equal to 0.13 bar and at oxygen partial

1032

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

pressures varying from 0.06 to 0.83 bar. The partial pressure sensitivity indicates increasing mass transfer and/or charge transfer resistance with decreasing oxygen concentration, c.f. the systematic investigation of the influence of ohmic resistance, charge transfer resistance and mass transfer resistance discussed by Yuh and Selman [26] and Lu et al. [27]. From the charge transfer resistance, it is possible to obtain formal reaction orders of oxygen and carbon dioxides from the dependence of the reciprocal formal charge transfer resistances (Rct =(dh/di )i = 0) on the partial pressures of carbon dioxide and oxygen. Fig. 5(a,b) shows such log(1/Rct) versus 1/T plots for the Li/K carbonate electrolyte at temperatures between 550 and 750°C and oxygen and carbon dioxide pres-

Fig. 5. (a) Logarithm of reciprocal charge transfer resistance vs. reciprocal temperature at various oxygen partial pressures with carbon dioxide partial pressure equal to 0.13 bar; (b) logarithm of reciprocal charge transfer resistance vs. reciprocal temperature for various carbon dioxide at constant partial oxygen pressures, pO2 = 0.13 bar. Electrolyte for both cases: Li/K2(CO3) eutectic.

sures varying from 0.04 to 0.8 bar with either pCO2 or pO2 maintained constant at 0.13 bar. It is obvious that the formal reaction orders are not constant, and change somewhat with partial pressures, and strongly with temperature. The latter is due to the change in activation energy at temperatures around 630 – 640°C. Below this temperature, the effective activation energy is almost double that observed above 630°C. The reaction order nO2 is lower at oxygen partial pressures B0.12 bar, than at oxygen partial pressures \0.12 bar. The difference between the two partial pressure ranges is also greater at higher temperatures, and vanishes at approximately 550°C, where the formal reaction order of oxygen measured from the charge transfer resistance nO2 is 0.5 9 0.1. It is somewhat \ 0.5 for high and low oxygen partial pressures. However, it diminishes at high temperatures to a value of 0.4 for pO2 \ 0.12 bar, and decreases steadily from 0.5 to a value of 0.3 at 750°C at low oxygen partial pressures (pO2 B 0.12 bar). In Fig. 7 the slightly different behaviour for low and high oxygen partial pressures is neglected to demonstrate the decrease from nO2 = 0.5 at low temperatures to 0.3 at high temperatures. The formal reaction order of carbon dioxide according to Fig. 5(b) is close to zero (even slightly negative) at low temperatures, but increases steadily from zero at 550°C to 0.5 at 750°C at low carbon dioxide pressures, as shown in Fig. 7(b). The activation energy changes from lower values of approximately 50 – 60 kJ mol − 1 at higher temperatures to higher values of around 90 – 100 kJ mol − 1 at lower temperatures. It is clear from Fig. 5(a) that also the activation energy is somewhat influenced by the oxygen and carbon dioxide partial pressures.

3.1.3.2. O2 reduction macrokinetics in the Li2CO3 / Na2CO3 eutectic. Fig. 4 also shows (dotted lines) two different current – voltage curves for cathodic oxygen reduction in laboratory fuel cells with Li/Na carbonate eutectic melts for high (0.87 bar) and low (0.06 bar) oxygen partial pressures under otherwise identical conditions. The difference of overpotential – current density correlations in the two electrolytes shows lower charge transfer resistance at high pO2 but definitely higher charge transfer resistance in Li/Na carbonate than in Li/K carbonate at lower oxygen partial pressures. At higher current densities, higher mass transfer resistances for oxygen at low partial pressures result in a higher dh/di slope and oxygen overpotential in Li/Na carbonate melt than in Li/K carbonate. Table 2 compares the charge transfer resistances observed in MCFC cathodes with Li/K and Li/Na carbonate eutectic melts at 650°C with dry cathode gas (1 bar total, with 0.13 bar carbon dioxide) and different oxygen partial pressures in the range from 0.06 to 0.87 bar (balance nitrogen). In both electrolytes the charge transfer resistance increases with

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

1033

Fig. 7. (a) Change of the formal reaction order of oxygen and (b) formal reaction order of carbon dioxide vs. temperature for Li/K and Li/Na eutectic melts. Fig. 6. (a) Logarithm of reciprocal charge transfer resistance vs. reciprocal temperature at various oxygen partial pressures with carbon dioxide partial pressure equal to 0.13 bar; (b) logarithm of reciprocal charge transfer resistance vs. reciprocal temperature for various carbon dioxide partial pressures with pO2 = 0.13 bar. Electrolyte for both cases: (Li/Na) eutectic.

decreasing oxygen pressure. At high oxygen partial pressures (0.87–0.15 bar) the cathodic charge transfer resistance is lower in Li/Na carbonate electrolyte than in Li/K carbonate eutectic. Below this partial pressure range, the charge transfer resistance increases according to a first order reaction (1/Rct pO2) in Li/Na carbonate, but as a 0.5 order process (1/Rct p 0.5 O2 ) in Li/K carbonate melt. Therefore, at 0.06 bar, the charge transfer resistance is by approximately a factor of two higher in Li/Na than in Li/K carbonate melts. The difference noted at small current densities becomes even more pronounced at higher current densities so that at 150 mA cm − 2 and 0.06 bar oxygen pressure the cathode overpotential in Li/K carbonate is − 200 mV, but is as high as −400 mV in Li/Na carbonate. As is shown in Fig. 7(a), the formal oxygen reaction order, which is close to 0.5 throughout the whole

temperature range in Li/K carbonate melts, is substantially higher in Li/Na melts. It is close to 1.0 at lower temperatures but decreases to a value of 0.7 at higher temperatures as a result of the change in activation energy around 630°C.

Table 2 Charge transfer resistance of MCFC oxygen cathodes in Li/K and Li/Na carbonate eutectic melts at 650°C with pCO2 =0.1 bar* pO2

(Li/K) carbonate Rct (V cm2)

(Li/Na) carbonate Rct (V cm2)

0.06 0.80 0.10 0.12 0.15 0.25 0.50 0.87

1.5 1.28 1.10 1.00 0.98 0.78 0.60 0.52

2.72 1.82 1.34 1.10 0.80 0.60 0.42 0.33

* Mean error of the resistance is B10 %.

1034

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

Fig. 8. Increase of cell resistance at vanishing and small current density on going from the inlet to the outlet of MCFC cathodes operated on lean cathode gas, as calculated for Li/K and Li/Na eutectic electrolytes.

The detailed dependence of formal reaction orders of carbon dioxide in Li/Na and Li/K carbonate melts on temperature is plotted in Fig. 7(b). In Li/Na carbonate the carbon dioxide pressure dependence of the charge transfer resistance for O2 reduction results in a reaction order close to but somewhat less than, unity at low temperatures, decreasing to approximately 0.7 at higher temperatures. These data allow calculation of how the effective total cell resistance of MCFCs operating on lean cathode gas changes at 650°C from cathode inlet to cathode outlet due to CO2 and O2 depletion from 0.1 and 0.13 bar, respectively, to 0.06 and 0.05 bar, respectively, in both melts operating on relatively lean cathode gas.1 Fig. 8 shows the increase in charge transfer resistance brought about by simultaneous depletion of oxygen and carbon dioxide along the cell. With Li/K carbonate it has a starting value of 1.3 V cm2, which remains almost constant along two thirds of the cell length. At the end, it becomes 2.5 V cm2. With Li/Na carbonate it rises from 1.7 to 3.7 V cm2. 4. Discussion With the exception of those concerning charge transfer resistances in MCFCs with Li/ Na carbonate elec1

The depletion of CO2 and O2 proceeds according to the stoichiometry of 2:1 and the inlet partial pressures of both gases are calculated according to a l-value of 1:1.5 for the whole MCFC process.

trolyte these investigations are not really new. The electrode kinetic measurements for oxygen reduction at immersed flag electrodes were performed to check previous and partially controversial results of other authors. Yuh and Selman [28] and also Tomczyk [29] have recently reviewed the methods used and results obtained by different authors. We will not further review previous work, but stress, that our results on the kinetics of the cathodic oxygen reduction obtained by pulse potentiometry at gold electrodes are very close to those obtained by Lu and Selman [22]. This may be seen from the exchange current density and reaction orders from oxygen and carbon dioxide at 650°C agreement within the given limits of uncertainty [22] in Table 3. The activation energy obtained (53 kJ mol − 1) is close to the low temperature value obtained by Uchida et al. (59 kJ mol − 1) [11]. Mass transfer reaction orders for dissolved oxygen species obtained from steady-state current – voltage curves with natural convection also agree quite well with the data of Yamada et al. [30]. The results of Selman and the present oxygen transport reaction orders (0.5 for oxygen and close to zero for carbon dioxide) do not agree with the values obtained from impedance spectra at equilibrium, i.e. by analysis of the s terms of the Warburg impedance [31]. A straightforward explanation for this discrepancy is not evident. However, we acknowledge Tomczyk’s approach [29], for reconciling the contradictory results, even though it is not completely satisfactory. Our data clearly support the peroxocarbonate mechanism (O2 + 2CO23 − “ 2CO24 − ) postulated by Tomczyk (the s value analyses of different authors indicate that peroxide (O22 − ) and

Table 3 Comparison of microkinetic data for cathodic oxygen reduction with data of Selman and co-workers obtained in (Li/ K)2CO3 eutectic at 650°C at 0.9 bar O2 and 0.1 bar CO2 Fast methods: reaction orders of charge transfer Fast pulse methods: electrode kinetics at gold electrodes: i0 (mA cm−2) Lu and Selman [22] This work 16.5 13.6 Reaction orders: nO2 0.26 0.33 90.1 nCO2 0.16 0.13 90.1 Slow methods: reaction orders of O2 transport Yamada et al. [30] This work nO2 0.49 0.47 9 0.04 nCO2 −0.11 0.05 9 0.05

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

1035

Table 4 Exchange current densities of oxygen reduction in Li/K carbonate eutectic at 650°C Electrode

Conditionsa

i0 (mA cm−2)

CD (mF cm−2)

Ref.

NiO powder Au/NiO d = 43–172 nm galv. depos. Au/NiO d = 43–172 nm galv. depos. NiO (monocrystal) LiCoO2 (powder) In-situ oxidised monocrystal

(a) (b) (b) (c) (a) (c)

86.4 16.7–19.1 14.4–19.3 17 12.4 220

140 90–154 75–110 50 131 190–320

[33] [34] [34] [18] [33] [19]

a

Conditions: (a) Li/K carbonate eutectic, 650°C, O2/CO2 = 0.9/0.1; (b) as under (a) but Li/K= 42.7/57.3 (mol/mol); (c) as under (a) but T = 727°C.

superoxide (O− 2 ) transport together with dissolved CO2 transport is mass transfer determining). Our activation energy for mass transport of dissolved oxygen species (57 kJ mol − 1) differs substantially from that reported by Uchida and co-workers [11] (25 kJ mol − 1 for low and 150 kJ mol − 1 for high temperatures). As we did not investigate the microkinetics of cathodic O2 reduction in more detail because we were mainly interested in the difference of macrokinetics in Li/K and Li/Na carbonate melts, we will not try to explain these differences. In practice, it is the macrokinetics of cathodic oxygen reduction which determine practical MCFC performance. It is sufficient to mention Yu and Selman’s earlier finding because O2 − neutralisation by CO2 is slow [32] and may indeed become rate determining. A consistent analysis of all relevant kinetic data obtained at different O2 and CO2 pressure and at different temperatures may not yet have been achieved by comparing available results. Claes et al. [33] have also indicated, that two types of dissolved carbon dioxide (physically dissolved and pyrocarbonate) must be taken into consideration. This may have further consequences in explaining O2 reduction kinetics. It is evident from the reaction orders of Table 3 that charge transfer kinetics in cathodic oxygen reduction and mass transport data of dissolved oxygen species are not inconsistent with Tomczyk’s assumption [29], that oxygen dissolves at least to an essential part in lithium/ potassium carbonate eutectic as peroxycarbonate according to Eq. (11). O2 +2CO23 − ? 2CO24 −

(11)

and that the cathodic reduction of dissolved oxygen is mainly due to peroxycarbonate reduction (Eq. (12)). CO24 − +2e− “CO23 − +O2 −

(12)

More important in our microkinetic investigation is the discrepancy between the present and previous results in the ratio of the exchange current densities on lithiated nickel oxide and lithium cobaltite. Our data for oxygen reduction on nickel oxide is substantially

lower, but not by an order of magnitude, than most other data reported in the literature. For instance, Hatoh et al. [34] reported i0 values of 86 mA cm − 2, an unusually high value, and Uchida et al. [35] reported values ranging from 10 to 20 mA cm − 2, which is the range of the most frequently reported values. It is obvious that the structure of the porous oxidic electrode determines the true surface exposed to the electrolyte and, therefore, the true and measured effective current density. Our method of extrapolating the apparent current densities to a coating thickness of 0.25 mm is the way to overcome this uncertainty in true surface and yet it is still ambiguous. Hence, the uncertainty in the current densities as extrapolated is perhaps not identical to the uncertainty of the extrapolation to a given thickness. The real uncertainty of the true and geometric surface areas may not be more precise than by a factor of two. That exchange current densities on LiCoO2 cathodes are higher than on lithiated nickel oxide had never been previously reported. For instance, Hatoh et al. reported values which were by a factor of seven lower for LiCoO3 than for NiO [34]. However, our result on LiCoO2 is well in the range of most reported values. Table 4 gives some further literature data and compares them with the respective surface specific electrode capacitances in order to show the scatter of the data and to demonstrate that there is no clear correlation between exchange current density and electrode capacitance. Therefore, we simply conclude that lithiated nickel oxide and lithium cobaltate are comparable in their electrocatalytic activity for oxygen reduction in carbonate melts, and that the choice of the cathode material is not likely to be a critical question for MCFC operation as O2 reduction is a relatively fast process at both materials. Macrokinetic and microkinetic data in Li/K carbonate eutectic are in relatively good agreement with each other in the low temperature range from 550 to 650°C, since the reaction orders are almost the same for microkinetic O2 mass transport and charge transfer and

1036

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037

for macrokinetic O2 reduction. The orders are in all cases close to 0.5 for oxygen and close to zero for carbon dioxide. The macrokinetic current–voltage curves in the last part of this investigation shows, that the nature of the electrolyte is of much higher importance for fuel cell performance than that of the oxide cathode material. There is a clear distinction between the two electrolytes studied in respect to high and low carbon dioxide and oxygen pressures in the cathode gas. Most frequently current–voltage curves have been reported for standard cathode gas compositions of 33/ 66 vol./vol. ratio oxygen to carbon dioxide. At such ‘fat’ gas compositions, an oxygen cathode at 650°C with Li/Na carbonate eutectic should have some kinetic advantage compared to Li/K carbonate eutectic. This is seen from the charge transfer resistances of Table 2. Kim and Selman [36] were the first to measure current– voltage curves in laboratory MCFCs with lean cathode gas, which reduces NiO in the melt. They found drastically increased overpotentials under these conditions, which represent the cathode gas compositions in real operating MCFCs. To prevent excessive balance-ofplant costs, today’s MCFC uses a very simple gas process technology in which the cathode gas is prepared by mixing combusted depleted anode gas with the amount of air corresponding to an overall oxygen to anode inlet fuel ratio (l ratio) of approximately 1.5, so that the inlet concentrations of oxygen and carbon dioxide do not exceed 10 and 13%, respectively. These are depleted to 6 and 5%, respectively (balance mainly N2 and some water vapour) at the cathode exit. Therefore, the current–voltage curves measured with these relatively ‘lean’ gases are most relevant in practice. With lean cathode gases, the Li/Na carbonate eutectic is now clearly at an disadvantage compared to Li/K carbonate melt. The effect of the formal charge transfer reaction order of close to unity for oxygen and carbon dioxide across the whole concentration range of both gases, which prevails in the cell increases the charge transfer resistance for cathodic oxygen reduction in Li/Na carbonate eutectic to very high values for lean gas. Fig. 8 shows the calculated change of the cell resistance (the sum of charge transfer resistance and ohmic cell resistance) in going from the cathode inlet to the outlet. Whereas in Li/K carbonate eutectic the cell resistance changes by 50%, it almost triples from the inlet to the outlet in Li/Na carbonate eutectic.

5. Conclusions From the above data, there appears to be little point in substituting Li/Na carbonate eutectic for Li/K eutectic in the MCFC unless changes in cathode design can be used to reduce mass transfer resistance in Li/Na carbonate eutectic. In this electrolyte the oxygen solu-

bility is lower by a factor of two, and the carbon dioxide partial pressure reaction order for charge transfer is higher, strongly increasing the charge transfer resistance for oxygen reduction compared with Li/K eutectic melt with lean cathode gases. However, this applies to non-pressurised cell operation, and does not hold for pressurised cells. Ishikawajima Harima Heavy Industries are developing MCFC systems pressurised in the range 5 – 7 atm. Under these conditions 5 vol.% would correspond to a partial pressure of 0.25 – 0.35 bar, so that charge transfer resistances with Li/Na eutectic melt would still be tolerable [37]. A final remark concerns the change in activation energy for O2 reduction in the cell from high values at low temperature to low values at high temperatures, as shown in Figs. 5 and 6. These data completely contradict Nishina, Uchida and Selman’s findings for oxygen transport towards smooth flag electrodes. This is not surprising, since the present results are for current – voltage curves at porous electrodes. The effect can be explained by the well-established influence of the temperature-dependent utilisation of porous catalysts on the effective activation energy of reactions catalysed on them [38]. As temperature and catalytic rates increase, mass transfer hindrance in the porous catalyst becomes more important. Above a certain threshold temperature the utilisation of catalyst particles is diminished sizeably below unity — as the so called Thiele modulus increases beyond a value of three [38]. Above this threshold temperature, the activation energy is halved. Therefore the effective phenomena in the present and Uhida and Selman’s measurements become different.

Acknowledgements The authors are indebted to MTU, Friedrichshafen AG and the German Ministry of Research and Development for financing essential parts of this investigation.

References [1] C.H.J. Broers, M. Schenke, in: G.J. Young (Ed.), Fuel Cells, vol. II, Reinhold, New York, 1963. [2] H. Mohn, H. Wendt, Z. Physik. Chem. 192 (1995) 101. [3] P. Tomczyk, J. Electroanal. Chem. 379 (1994) 335. [4] H. Urushibata, T. Murahashi, Proceedings of the International Fuel Cell Conference (IFCC), Makuhari, Japan, 1992, p. 223. [5] M. Cassir, M. Ohivri, V. Albin, B. Malinowska, J. Devynck, J. Electroanal. Chem. 452 (1998) 127. [6] J.R. Selman, H.C. Maru, Physical Chemistry and Electrochemistry of Alkali Carbonate Melts — Advances in Molten Salt Chemistry, vol. 4, Plenum Press, New York, 1981.

K. Janowitz et al. / Electrochimica Acta 45 (1999) 1025–1037 [7] P. Delahay, J. Phys. Chem. 66 (1962) 2204. [8] H.P. van Leeuwen, Electrochim. Acta 23 (1978) 207. [9] I. Uchida, T. Nishina, Y. Nugikura, K. Itaya, J. Electroanal. Chem. 206 (1986) 127. [10] I. Uchida, Y. Nugikura, T. Nishina, K. Itaya, J. Electroanal. Chem. 206 (1986) 229. [11] I. Uchida, T. Nishina, Y. Nugikura, K. Itaya, J. Electroanal. Chem. 209 (1986) 125. [12] A.J. Appleby, S.B. Nicholson, J. Electroanal. Chem. 53 (1974) 105. [13] A.J. Appleby, S.B. Nicholson, J. Electroanal. Chem. 83 (1977) 309. [14] A.J. Appleby, S.B. Nicholson, J. Electroanal. Chem. 112 (1980) 71. [15] S.H. White, U.M. Twardoch, J. Electrochem. Soc. 135 (1988) 892. [16] M. Matsumura, J.R. Selman, J. Electrochem. Soc. 139 (1992) 1255. [17] Ullmanns’ Encyclopa¨die der Technischen Chemie, vol. A9, Verlag Chemie, Weinheim, 1991, p. 151. [18] P. Tomczyk, H. Sato, K. Yamada, T. Nishina, I. Uchida, J. Electroanal. Chem. 391 (1995) 125. [19] P. Tomczyk, H. Sato, K. Yamada, T. Nishina, I. Uchida, J. Electroanal. Chem. 391 (1995) 133. [20] J.R. Selman, L.G. Marianowski, Fuel cells, in: D.G. Lovering (Ed.), Molten Salt Technology, Plenum Press, New York, 1982. [21] O. Bo¨hme, F.U. Leidich, H.J. Salge, H. Wendt, Int. J. Hydrogen Energy 19 (1994) 349. [22] S.-H. Lu, J.R. Selman, J. Electroanal. Chem. 333 (1992) 257.

1037

[23] B.B. Dave, R.E. White, S. Srinivasan, A.J. Appleby, J. Electrochem. Soc. 138 (1991) 673. [24] B.B. Dave, R.E. White, S. Srinivasan, A.J. Appleby, J. Electrochem. Soc. 138 (1991) 2657. [25] B.B. Dave, R.E. White, S. Srinivasan, A.J. Appleby, J. Electrochem. Soc. 140 (1993) 2139. [26] C.Y. Yuh, J.R. Selman, J. Electrochem. Soc. 131 (1984) 2062. [27] G.L. Lu, J.R. Selman, L. Pomp, J. Electrochem. Soc. 140 (1993) 390. [28] C.Y. Yuh, J.R. Selman, J. Electrochem. Soc. 138 (1991) 3649. [29] P. Tomczyk, J. Electroanal. Chem. 379 (1994) 353. [30] K. Yamada, T. Nishina, I. Uchida, J.R. Selman, Electrochim. Acta 38 (1993) 2405. [31] K. Yamada, T. Nishina, I. Uchida, Electrochim. Acta 40 (1995) 1927. [32] S.H. Lu, J.R. Selman, J. Electrochem. Soc. 137 (1990) 1125. [33] P. Claes, B. Thirion, J. Gilbert, Electrochim. Acta 41 (1996) 141. [34] K. Hatoh, U. Nikura, E. Yasumoto, T. Gamo, J. Electrochem. Soc. 141 (1944) 1725. [35] I. Uchida, Y. Mugikura, T. Nishina, K. Itaya, J. Electroanal. Chem. 206 (1986) 241. [36] Y.H. Kim, J.R. Selman, in: C. Hussey, D.S,. Newman, Y. Ito (Eds.), Molten Salts, Proc. vol. PV 94-13, The Electrochemical Society, Pennington, 1994, NJ, p. 781. [37] Communicated by I. Toshiba at the MCFC Workshop, Delft, 8th January, 1999. [38] R. Aris, Chem. Eng. Sci. 6 (1957) 262.